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This 6 page Class Notes was uploaded by PAULA CANGOMA on Tuesday April 12, 2016. The Class Notes belongs to math and chemestry at University of North Dakota taught by David Pierce in Spring 2016. Since its upload, it has received 25 views. For similar materials see General Chemistry I in Engineering and Tech at University of North Dakota.
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Date Created: 04/12/16
Chemical Tools: Experimentation and Measurement • Recognize the seven basic SI units of measure. • Give the numerical equivalent of the common metric prefixes used with these units. • Express numbers in scientific notation. • Determine the number of significant digits in a measured quantity. • State the result of a calculation involving measured quantities to the correct number of significant digits. • Convert temperatures between the Celsius and Kelvin scales. • Relate density to mass and volume. • Set up and solve problems using dimensional analysis. CHAPTER 1 – The Structure and Stability of Atoms • Give the symbol and name of the first 36 elements of the periodic table. • Identify the group and period to which an element belongs. • Identify the regions of the periodic table (metals, metalloids, nonmetals, alkali metals, alkaline earth metals, transition metals). • Identify properties as physical or chemical. • Identify physical and chemical properties as either extensive or intensive. • Explain the laws of definite proportions, multiple proportions, and conservation of mass. • Show that the law of multiple proportions is obeyed for two different compounds comprised of the same two elements. • Summarize Dalton’s Atomic Theory. • Explain what information about atomic structure was revealed by the experiments of Thomson, Millikan, and Rutherford. • Describe the atom in terms of composition, mass, and volume of the nucleus relative to the mass and volume occupied by the electrons. • Given the symbol for an isotope of an atom or ion, determine the number of protons, neutrons, and electrons. • Given the mass and natural abundance of all isotopes of a given element, calculate the average atomic mass of that element. • For any element, calculate the mass in grams of a single atom and the number of atoms in a given number of grams. • Calculate molar mass of an element. CHAPTER 2 – Periodicity & the Electronic Structure of Atoms • Relate wavelength, frequency, and energy of electromagnetic radiation. • Understand the clues that atomic line spectra provide to the electronic structure of atoms. • Use a set of quantum numbers to describe a particular orbital. • Sketch and name each of the s, p, and d orbitals. • Predict ground-state electron configurations for the first 36 elements (through Kr). • Use orbital-filling diagrams to determine the number of unpaired electrons in a given electron configuration, including the effects of Hund’s rule. • Understand the concept of valence electron configuration and its relationship to the structure of the periodic table. • Write a general valence-shell electron configuration for each group of the periodic table. • Know the major subdivisions (“blocks”) of the periodic table. • Identify the blocks in which the elements are located. • Given a set of atoms, determine which atom is expected to have the largest radius. CHAPTER 3 – Atoms and Ionic Bonds • Identify which substances are ionic and which are molecular. • Generate the appropriate formula for ionic compounds formed from simple monatomic ions. • Give formulas and names of common polyatomic ions and ionic compounds. • Predict ground‐state electron configurations for ions. • Predict the most likely ion formed from a given representative element. • Give the noble gas configuration of common cations and anions. • Given a set of ions, determine which ion is expected to have the largest radius. • Predict which has the higher first ionization energy for any pair of elements. • Predict the higher second, third, fourth, etc. ionization energy for any pair of elements. • Predict which has the larger electron affinity for any pair of elements. • Explain the Octet rule. • Describe ionic bonding and the formation of binary ionic compounds from their elements. • Predict which of two ionic compounds should have the greater lattice energy on the basis of ionic charges and ionic radii. • Use the valence electron configuration to generalize the chemistry of each family of elements studied in this chapter. • Give the formulas of products formed when alkali metals and alkaline earth metals react with halogens, hydrogen, oxygen, water, and ammonia. Balance the equations. • Give the formulas of products formed when halogens react with metals and hydrogen CHAPTER 4 – Atoms and Covalent Bonds • Describe covalent bonding and the formation of binary molecular compounds from their elements. • Use a table of electronegativity to predict which of two bonds is expected to be more polar. • Use the periodic table to predict which of two elements is more electronegative and whether a bond between them would be is ionic, polar covalent, or non‐polar covalent. • Properly name binary (two element) molecular compounds • Name and recognize by structural or condensed formulas simple alkane hydrocarbons. • Write electron-dot (Lewis) symbols for atoms and tell how many electrons must be shared to enable the atom to achieve a completed valence shell. • Give the symbol of the noble gas with the same number of valence electrons. • For each atom in an electron‐dot structure, give the number of bonded electron pairs and the number of non‐bonded electron pairs. • For a given electron-dot structure, give the number of single bonds, double bonds, and triple bonds. Give the bond order of each bond. • Draw electron-dot structures of molecules and polyatomic ions, employing multiple bonding and resonance structures as needed. • Determine the formal charge on each atom in a Lewis structure and use the formal charges to select the best Lewis structure. • Know when to expect the octet rule to be valid and when it can fails CHAPTER 5 – Covalent Bonds and Molecular Structure • Use the VSEPR model to predict geometries of molecules and polyatomic ions, including those with more than one central atom. • Identify common hybridization schemes for s and p orbitals and their correlation with molecular shapes. • Sketch and identify the orbitals used by each atom to form bonds in molecules and polyatomic ions. Show which orbital overlaps result in sigma (σ) bonds and which result in pi (π) bonds. • Know the composition of double and triple bonds in terms of π and σ components. CHAPTER 6 – Chemical Arithmetic: Stoichiometry • Write and balance chemical equations for simple chemical reactions. • Calculate molecular mass (or formula mass) of a compound. • Interconvert grams, moles, and numbers of formula units. • Determine the number of moles and grams of one reactant needed to react with a given number of moles and grams of another reactant, and the number of moles and grams of product(s) that result from the reaction. • Calculate percent yield. • Identify the limiting and excess reactants in a reaction mixture. • Determine the mass of excess reactant remaining at the end of a reaction and the mass of product(s) produced. • Calculate the mass of products produced from a given mass of reactants when the theoretical yield is < 100%. • Differentiate between molecular and empirical formulas. • Determine the molecular formula of a compound from empirical formula and molar mass. • Determine the percent composition and empirical formula of a compound. • Describe how to prepare a solution of known molarity by dissolving a solid in a solvent, and by diluting a more concentrated solution. • Interconvert solution molarity, solution volume, solute moles, and solute grams. • Determine the volume of one reactant solution needed to react with a given volume of a second reactant solution. • Determine the concentration of a species using data from an acid‐base titration or a redox titration CHAPTER 7 – Reactions in Aqueous Solutions • Identify common types of chemical reactions; specifically, precipitations reactions, acid-base neutralization reactions, and oxidation-reduction reactions. • Classify substances as electrolytes or nonelectrolytes. • Write molecular, ionic, and net ionic equations for precipitation, acid-base, and redox reactions. • Use solubility rules to predict whether a precipitate might form when aqueous salt solutions are mixed. • Identify which substances are acids and which are bases, and properly name common acids. • Identify common strong acids and strong bases. • Identify common polyprotic acids. • Recognize acid-base neutralization reactions. • Assign oxidation numbers to each atom in a chemical species. • In a redox reaction, identify the species oxidized, the species reduced, the oxidizing agent, and the reducing agent. • Using an activity series, predict whether a redox reaction will occur when a metal is placed in contact with a solution containing an ion of a different metal. • Balance simple redox reactions. CHAPTER 8 – Thermochemistry: Chemical Energy • Differentiate between the concepts of heat and temperature. • Differentiate between potential and kinetic energy. • Characterize a process as endothermic or exothermic. • Explain the first law of thermodynamics (energy is conserved). • Identify and describe features of a state function, like internal energy. • Recognize how energy is gained or lost by a system (like a chemical reaction or phase change) as heat (enthalpy change, ∆H). • Use the sign of ∆H to identify whether heat is being gained or lost by a system (like a chemical reaction or phase change). • Given a balanced chemical equation and enthalpy change for a chemical reaction, calculate the enthalpy change per mole or per gram of each reactant and product. • Perform calorimetry calculations involving specific heat (or molar heat capacity), heat flow, and temperature change. • Use Hess's law to determine ∆H values. • Use standard enthalpies of formation to calculate a standard enthalpy of reaction. • Use bond dissociation energies to calculate a standard enthalpy of a chemical reaction CHAPTER 9 – Gases: Their Properties and behavior • Know which elements exist as gases at 25°C and 1 atm pressure. • Know the four common physical characteristics of all gases. • Interconvert units of pressure. • Describe how the pressure of a gas is measured using a manometer. • Use the ideal gas law to calculate final pressure, volume, moles of gas, or temperature from initial pressure, volume, moles of gas, and temperature. • Use the ideal gas law to calculate pressure, volume, moles of gas, or temperature, given the other three variables. • Use the ideal gas law to determine the density of a gas. • Use the ideal gas law to determine the molar mass of a gas. • Perform stoichiometric calculations relating the mass of a reactant to the mass, moles, and volume or pressure of a gaseous product. • Use Dalton's law to calculate the partial pressure of a gas in a mixture. • Explain each of the gas laws using the Kinetic Molecular Theory. • Use Graham's law to calculate the relative rates of effusion of two different gases. • State conditions under which a gas is expected to behave ideally or non‐ideally. CHAPTER 10 – Liquids, Solids, and Phase Changes • Compare and contrast characteristic properties of gases, liquids, and solids. • Determine whether a molecule is expected to be polar or non-polar using its VSEPR geometry and electronegativity. • Identify major intermolecular forces present in substances and determine which of two substances will exhibit the stronger intermolecular force. • Identify molecules that can form hydrogen bonds. • Know and explain basic anomalous properties of water compared to its chemical analogs. • Describe how viscosity and surface tension vary with temperature and the strength of intermolecular forces. • Describe the different phase changes and explain their respective enthalpy changes. • Interpret the heating curve of a substance. • Understand the relationship between the rate of evaporation and vapor pressure of a substance. • Predict the relative boiling points, vapor pressures, and enthalpies of vaporization of two substances from the strengths of their intermolecular forces. • Distinguish metals, ionic solids, network solids, and molecular solids by their structures and by their properties. • Sketch a phase diagram, labeling the axes and each of the regions, and locate the triple point, critical point, normal melting point, and the normal boiling point. • Predict the relative densities of the liquid and solid phases of a substance from its phase diagram. • Identify the stable phase or phases of a substance at a given temperature and pressure from its phase diagram.
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